PACIFIC SOUTHWEST ASSOCIATION O F CHEMISTRY TEACHERS
A CHEMISTRY COURSE FOR HIGH-SCHOOL TEACHERS Shell Merit Fellowship Program RICHARD H. EASTMAN Stanford University, Palo Alto, California
DURING the Summer Quarter of 1956-57 I gave a course in general chemistry to fourteen high-school science teachers selected from the western half of the United States under the Shell Merit Fellowship Program a t Stanford University, where I regularly teach general chemistry and do research in organic chemistry. It is believed that some observations on the content and conduct of the course may prove useful to those engaged in similar activities in the future. The point of view adopted as far as content is concerned recognized that research activity which requires a university scientist to keep in touch with the developing frontier of his subject unfortunately has no similarly stimulating counterpart in the life of the high-school science teacher, no matter what the depth of his devotion to his subject. Therefore, the course was conceived by the lecturer as one whose primary purpose was to clarify the understanding of fundamental chemical concepts and bring the high-school teachers up-to-date by presenting those aspects of chemical theory which have developed rapidly in the last ten years. To secure t,his purpose in part, the course was limited in enrollment to the homogeneous group of high-school teachers, it being felt that inclusion of graduate or other students who desired a review of general chemistry for personal purposes would put a check-rein by their very presence on the free discussion which was hoped would and did in fact develop. To select the content of the course of forty lecturers, the instructor asked the students what chemical subjects they most desired to hear discussed. While this selection was being made, he himself, drew up a list of topics which he conceived to be important. That there was almost a one-to-one correspondence between the students' private choices and those of the instructor, could be interpreted in many ways. I n any event it certainly made the summer's work fun from everyone's point of view. That is not to say that the summer was an easy one, for, although I can speak only for the lecturer, it was without a doubt the most difficult job of teaching that he has ever done, and I suspect that the group enthusiasm exacted great effort from the students as well. But why not? Put fifteen highly intelligent people together who are interested in the same subject and what would anyone
expect? The subject exacted from all of us, everything of which we were capable. From the lecturer's point of view, it was most stimulating to sell his concept of his favorite subject to an audience of intellectual sponges who wanted most desperately to absorb and digest everything they were introduced to. But the flow was not all one way by any means. The lecturer got much in the way of understanding the problems of students from the high-school teachers. After all, the high-school student in chemistry is not far removed in maturity from the college freshman, and the lecturer feels sure that his teaching has benefited from the group experience. The esprit de corps and the enthusiasm which made possible this exchange would have been killed by the presence of outsiders whose special competence in other areas would have destroyed the atmosphere of free discussion. What, then, was discussed? First, the Law of Definite Composition and the nature of "pure" substances in the modern context. Naturally from this topic came a discussion of fundamental particles and isotopes, and thence, the development of a clear understanding of the difference between the Law of Conservation of Matter (weight of reactants = weight of products) and the Law of Conservation of Mass (matter mass energy mass of reactants = matter mass energy mass of products) and E = mc2. Useful in scaling the energy quantities in the later part of this discussion was the definition of energy quantities all the way from van der Waal's intermolecular forces to the magnitude of the forces bonding nucleons and the energies released in nuclear reactions. With the scale of energies defined, radioactivity was treated briefly from the point of view of the processes involved and the corresponding changes in mass and atomic number; and then, attention was directed to the chemical forces. The discussion of the types of chemical bonds was in terms of the two extremes- ionic and covalentand was preceded by development of the electron configurations of the atoms in their lowest energy states, using the mbital as the fundamental building block out of which were constructed suh-shells and principal shells following in a non-mathematical way C. A. Coulson ("Valence," Oxford University Press,
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JOURNAL OF CHEMICAL EDUCATION
London, 1952). Ionic bonds were regarded as being formed by electron transfer between atomic orbitals, and covalent bonds were treated from the molecular orbital or "polynuclear atom" point of view using linear combination of atomic orbitals and hybridization to arrive a t the qualitative geometrical aspects of covalent binding. Polar covalence was introduced using electronegativity, ionization potential and electron affinities t o rationalize the existence of dipoles. Secondary forces were discussed under van der Waal's and dipoledipole interaction. With the polar covalent bond concept well in hand, attention was turned to the destrnction and formation of covalent bonds in the reactions of proton acids with bases. The treatment of this subject was in terms of the Bronsted-Lowry General Acid-Base Theory and in addition to discussion of proton-transfer reactions in water solutions, the application of the theory to the non-aqueous system, concentrated sulfuric acid was developed following L. P. Hammett (''Physical Organic Chemistry," McGraw-Hill, New York, 1940, Chapt. 11), and using such examples as (CZHJZO H2S01
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(C2H&0H HSOa-, citing freezing-point lowering results to support the theory. As a prelude to the discussion of acids and bases, clear distinction was made between component and species concentration, and the part played by the latter in equilibrium constant expressions was developed. The quarter was brought to a close with application of chemical equilibrium concepts to oxidation reduction reactions, electrode processes and batteries. Aside from the formal lectures, the students were on several occasions taken through the chemistry department research laboratories to see the means by which modern chemical investigations are prosecuted. These tours were timed so that the particular processes studied in operation corresponded to the material under discussion in lecture at the time. Thus, when the definition of "pure substance" was under lecture discussion, the laboratory examination revealed fractionation and chromatographic columns, the apparatus for liquid-vapor partition chromatography, and the ultracentrifuge. With radioactivity came methods of counting radioactive decay events and the use of tracers in organic reactions, and when covalent bonds were discussed, recourse was had to the machines used to measure the absorption of ultraviolet and infrared radiation by molecules. It is hoped that this brief review of what went on a t Stanford in the chemistry course under the Shell Merit Fellowship program will be of service to instructors engaged in constructing similar courses. Over-all, the treatment was qualitative rather than quantitative, and aimed a t engendering an understanding of the present state of the science of chemistry in terms of broad concepts and their relationships. The text was used as a reference book, and was "Principles of Chemistry" by Hildebrand and Powell combined volume with "Reference Book of Inorganic Chemistry," by Latimer and Hildebrand [The Macmillan Company, New York, 1952.1 A topical outline of the subjects covered in the forty lectures during the eight weeks follows.
VOLUME 34, NO. 1, JLILY, 1957
Pure Substances Classificationof Systems Heterogenous, homogenous Accuracy and Precision Physical numbers, arithmetical numbers 11. Forces Magnitudes of Binding Forces Van der Waal forces. chemical bonds. intra-nuclear 111. TheLaw of Conservation of Mass Phlogiston Theory Energy-Mass Equivalence IV. Fundamental Particles Properties of Cathode Rays Quantitative Properties of Electrons e l m , e,
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Canal Rays Properties, hydrogen as residud gas, helium as residual gas Prout'sHypothesis Mass Defect Characterization of Atoms Determination of atomic number, mass number, isotopes, atomic mass, the mass speotrograph Radioactivity The Geiger counter, radioactive decay series, rule for stability of nuclei, synthetic isotopes, K-electron capture Electron Cloud (from atomic spectra) Postulates of quantum theory, energy levels of hydrogen atom, Heisenberg uncertainty principle, development of Schr6dinger equation, application to hydrogen and heavier atoms, energy sequence of subshells, electron configuration of the elements, Pauli's exclusion principle, Hund's rule. V. Cbemied Bands Examples of No Bond Ionic Bonds Covalent Bonds
EP, S p z , S P s
Oxidation States Variable "valence" Dipoles Effect of symmetry, dipole moment, dipole-dipole interaction Hydrogen Banding Dimerimtion, inter-molecular and intra-maleoulsr bonding VI. Acids and Bases
mole fraction Temperaturedependent methods: grams solute per liter of solution, molarity or formality, normality Electrolytes Strong, weak; molar concentration of water, autoprotolysis of water Lowry-Bronsted Concept of Acids Proton donors and acceptors, conjugate acids and bases Equilibrium Constants Dissociation, hydrolysis, ion product of water Titrstions Buffer Solutions Indioators End points and equivalence points, selection of indicators. measurine oH with indicators VII. Electrochemistry Fsreday's Laws Displacement Series Quantitative Measure of Strength of Redox Agents Half-Cell Reactions Factors Influencing Electrode Potentials and Processes Batteries
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